Hydroxide
Hydroxide is a diatomic anion with chemical formula −. It consists of an oxygen and a hydrogen atom held together by a covalent bond, and carries a negative electric charge. It is an important but usually minor constituent of water. It functions as a base, a ligand, a nucleophile and a catalyst. The hydroxide ion forms salts, some of which dissociate in aqueous solution, liberating solvated hydroxide ions. Sodium hydroxide is a multi-million-ton per annum commodity chemical. A hydroxide attached to a strongly electropositive center may itself ionize, liberating a hydrogen cation (H+), making the parent compound an acid. The corresponding electrically neutral compound •HO is the hydroxyl radical. The corresponding covalently-bound group -OH of atoms is the hydroxyl group. Hydroxide ion and hydroxyl group are nucleophiles and can act as a catalyst in organic chemistry. Many inorganic substances which bear the word "hydroxide" in their names are not ionic compounds of the hydroxide ion, but covalent compounds which contain hydroxyl groups. Hydroxide ion The hydroxide ion is a natural part of water, because of the self-ionization reaction: : H+ + OH− The equilibrium constant for this reaction, defined as :Kw = H+OH−H+ denotes the concentration of hydrogen cations and OH− the concentration of hydroxide ions has a value close to 10−14 at 25 °C, so the concentration of hydroxide ions in pure water is close to 10−7 mol∙dm−3, in order to satisfy the equal charge constraint. The pH of a solution is equal to the decimal cologarithm of the hydrogen cation concentration;Strictly speaking pH is the cologarithm of the hydrogen cation activity the pH of pure water is close to 7 at ambient temperatures. The concentration of hydroxide ions can be expressed in terms of pOH, which is close to 14 − pH,p(OH) signifies the minus the logarithm to base 10 of {OH−}, alternatively the logarithm of 1/{OH−} so pOH of pure water is also close to 7. Addition of a base to water will reduce the hydrogen cation concentration and therefore increase the hydroxide ion concentration (increase pH, decrease pOH) even if the base does not itself contain hydroxide. For example, ammonia solutions have a pH greater than 7 due to the reaction + + NH4+, which results in a decrease in hydrogen cation concentration and an increase in hydroxide ion concentration. pOH can be kept at a nearly constant value with various buffer solutions. In aqueous solution the hydroxide ion is a base in the Brønsted–Lowry sense as it can accept a protonIn this context proton is the term used for a solvated hydrogen cation from a Brønsted–Lowry acid to form a water molecule. It can also act as a Lewis base by donating a pair of electrons to a Lewis acid. In aqueous solution both hydrogen and hydroxide ions are strongly solvated, with hydrogen bonds between oxygen and hydrogen atoms. Indeed, the bihydroxide ion H3O2− has been characterized in the solid state. This compound is centrosymmetric and has a very short hydrogen bond (114.5 pm) that is similar to the length in the bifluoride ion HF2− (114 pm). In aqueous solution the hydroxide ion forms strong hydrogen bonds with water molecules. A consequence of this is that concentrated solutions of sodium hydroxide have high viscosity due to the formation of an extended network of hydrogen bonds as in hydrogen fluoride solutions. In solution, exposed to air, the hydroxide ion reacts rapidly with atmospheric carbon dioxide, acting as an acid, to form, initially, the bicarbonate ion. :OH− + HCO3− The equilibrium constant for this reaction can be specified either as a reaction with dissolved carbon dioxide or as a reaction with carbon dioxide gas (see carbonic acid for values and details). At neutral or acid pH, the reaction is slow, but is catalyzed by the enzyme carbonic anhydrase, which effectively creates hydroxide ions at the active site. Solutions containing the hydroxide ion attack glass. In this case, the silicates in glass are acting as acids. Basic hydroxides, whether solids or in solution, are stored in air-tight plastic containers. The hydroxide ion can function as a typical electron-pair donor ligand, forming such complexes as Al(OH)4−. It is also often found in mixed-ligand complexes of the type MLx(OH)yz+, where L is a ligand. The hydroxide ion often serves as a bridging ligand, donating one pair of electrons to each of the atoms being bridged. As illustrated by Pb2(OH)3+, metal hydroxides are often written in a simplified format. It can even act as a 3 electron-pair donor, as in the tetramer PtMe3OH4).Greenwood, p. 1168 When bound to a strongly electron-withdrawing metal centre, hydroxide ligands tend to ionises into oxide ligands. For example, the bichromate ion HCrO4− dissociates according to :O3CrO-H− CrO42− + H+ with a pKa of about 5.9.IUPAC SC-Database A comprehensive database of published data on equilibrium constants of metal complexes and ligands Vibrational spectra The infrared spectra of compounds containing the OH functional group have strong absorption bands in the region centered around 3500 cm−1. The high frequency of molecular vibration is a consequence of the small mass of the hydrogen atom as compared to the mass of the oxygen atom and this makes detection of hydroxyl groups by infrared spectroscopy relatively easy. A band due to an OH group tends to be sharp. However, the band width increases when the OH group is involved in hydrogen bonding. A water molecule has an HOH bending mode at about 1600 cm−1, so the absence of this band can be used to distinguish an OH group from a water molecule. When the OH group is bound to a metal ion in a coordination complex, an M−OH bending mode can be observed. For example, in Sn(OH)62− it occurs at 1065 cm−1. The bending mode for a bridging hydroxide tends to be at a lower frequency as in [(bipyridine)Cu(OH)2Cu(bipyridine)]2+ (955 cm−1).Nakamoto, Part B, p. 57 M−OH stretching vibrations occur below about 600 cm−1. For example, the tetrahedral ion Zn(OH)42− has bands at 470 cm−1 (Raman-active, polarized) and 420 cm−1 (infrared). The same ion has an (OH)Zn(OH) bending vibration at 300 cm−1. Chapter 5. Applications Sodium hydroxide solutions, also known as lye and caustic soda, are used in the manufacture of pulp and paper, textiles, drinking water, soaps and detergents, and as a drain cleaner. Worldwide production in 2004 was approximately 60 million tonnes. The principal method of manufacture is the chlor-alkali process. Solutions containing the hydroxide ion are generated when a salt of a weak acid is dissolved in water. Sodium carbonate is used as an alkali, for example, by virtue of the hydrolysis reaction :CO32− + H2O HCO3− + OH−; (pKa2 = 10.33 at 25 °C and zero ionic strength) Although the base strength of sodium carbonate solutions is lower than a concentrated sodium hydroxide solution, it has the advantage of being a solid. It is also manufactured on a vast scale (42 million tonnes in 2005) by the Solvay process.Kostick, Dennis (2006). "Soda Ash", chapter in 2005 Minerals Yearbook, United States Geological Survey. An example of the use of sodium carbonate as an alkali is when washing soda (another name for sodium carbonate) acts on insoluble esters, such as triglycerides, commonly known as fats, to hydrolyze them and make them soluble. Bauxite, a basic hydroxide of aluminium, is the principal ore from which the metal is manufactured. Similarly, goethite (α-FeO(OH)) and lepidocrocite (γ-FeO(OH)), basic hydroxides of iron, are among the principal ores used for the manufacture of metallic iron. Numerous other uses can be found in the articles on individual hydroxides. Inorganic hydroxides Alkali metals Aside from NaOH and KOH, which enjoy very large scale applications, the hydroxides of the other alkali metals also are useful. Lithium hydroxide is a strong base, with a pKb of -0.36.Lew. Kristi., Acids and Bases (Essential Chemistry). Infobase Publishing (2009). p43. Lithium hydroxide is used in breathing gas purification systems for spacecraft, submarines, and rebreathers to remove carbon dioxide from exhaled gas. :2 LiOH + CO2 → Li2CO3 + H2O The hydroxide of lithium is preferred to that of sodium because of its lower mass. Sodium hydroxide, potassium hydroxide and the hydroxides of the other alkali metals are also strong bases.Holleman, p. 1108 Alkaline earth metals Beryllium hydroxide Be(OH)2 is amphoteric.Thomas R. Dulski A manual for the chemical analysis of metals, ASTM International, 1996, ISBN 0-8031-2066-4 p. 100 The hydroxide itself is insoluble in water, with a solubility product log K*sp of −11.7. Addition of acid gives soluble hydrolysis products, including the trimeric ion Be3(OH)3(H2O)63+, which has OH groups bridging between pairs of beryllium ions making a 6-membered ring. At very low pH the aqua ion Be(H2O)42+ is formed. Addition of hydroxide to Be(OH)2 gives the soluble tetrahydroxo anion Be(OH)42−. The solubility in water of the other hydroxides in this group increases with increasing atomic number.Housecroft, p. 241 Magnesium hydroxide Mg(OH)2 is a strong base as are the hydroxides of the heavier alkaline earths, calcium hydroxide, strontium hydroxide and barium hydroxide. A solution/suspension of calcium hydroxide is known as limewater and can be used to test for the weak acid carbon dioxide. The reaction Ca(OH)2 + CO2 Ca2+ + HCO3− + OH− illustrates the strong basicity of calcium hydroxide. Soda lime, which is a mixture of NaOH and Ca(OH)2, is used as a CO2 absorbent. Boron group elements The simplest hydroxide of boron B(OH)3, known as boric acid, is an acid. Unlike the hydroxides of the alkali and alkaline earth hydroxides, it does not dissociate in aqueous solution. Instead, it reacts with water molecules acting as a Lewis acid, releasing protons. :B(OH)3 + H2O [[B(OH)4−]] + H+ A variety of oxyanions of boron are known, which, in the protonated form, contain hydroxide groups.Housectroft, p. 263 Aluminium hydroxide Al(OH)3 is amphoteric and dissolves in alkaline solution. :Al(OH)3 (solid) + OH− (aq) [[Al(OH)4−]] (aq) In the Bayer processBayer process chemistry for the production of pure aluminium oxide from bauxite minerals this equilibrium is manipulated by careful control of temperature and alkali concentration. In the first phase, aluminium dissolves in hot alkaline solution as Al(OH)4− but other hydroxides usually present in the mineral, such as iron hydroxides, do not dissolve because they are not amphoteric. After removal of the insolubles, the so-called red mud, pure aluminium hydroxide is made to precipitate by reducing the temperature and adding water to the extract, which, by diluting the alkali, lowers the pH of the solution. Basic aluminium hydroxide (OH), which may be present in bauxite, is also amphoteric. In mildly acidic solutions the hydroxo complexes formed by aluminium are somewhat different from those of boron, reflecting the greater size of Al(III) vs. B(III). The concentration of the species Al13(OH)327+ is very dependent on the total aluminium concentration. Various other hydroxo complexes are found in crystalline compounds. Perhaps the most important is the basic hydroxide AlO(OH), a polymeric material known by the names of the mineral forms boehmite or diaspore, depending on crystal structure. Gallium hydroxide, indium hydroxide and thallium(III) hydroxides are also amphoteric. Thallium(I) hydroxide is a strong base.James E. House Inorganic chemistry, Academic Press, 2008, ISBN 0-12-356786-6, p. 764 Carbon group elements Carbon forms no simple hydroxides. The hypothetical compound C(OH)4 is unstable in aqueous solution: :C(OH)4 → HCO3− + H3O+ :HCO3− + H+ H2CO3 Carbon dioxide is also known as carbonic anhydride, meaning that it forms by dehydration of carbonic acid H2CO3 (OC(OH)2).Greenwood, p. 310 Silicic acid is the name given to a variety of compounds with a generic formula SiOx(OH)4−2xn.Greenwood, p. 346R. K. Iler, The Chemistry of Silica, Wiley, New York, 1979 ISBN 0-471-02404-X Orthosilicic acid has been identified in very dilute aqueous solution. It is a weak acid with pKa1 = 9.84, pKa2 = 13.2 at 25 °C. It is usually written as H4SiO4 but the formula (OH)2 is generally accepted . Other silicic acids such as metasilicic acid (H2SiO3), disilicic acid (H2Si2O5), and pyrosilicic acid (H6Si2O7) have been characterized. These acids also have hydroxide groups attached to the silicon; the formulas suggest that these acids are protonated forms of polyoxyanions. Few hydroxo complexes of germanium have been characterized. Tin(II) hydroxide Sn(OH)2 was prepared in anhydrous media. When tin(II) oxide is treated with alkali the pyramidal hydroxo complex Sn(OH)3− is formed. When solutions containing this ion are acidified the ion Sn3(OH)42+ is formed together with some basic hydroxo complexes. The structure of Sn3(OH)42+ has a triangle of tin atoms connected by bridging hydroxide groups.Greenwood, p. 384 Tin(IV) hydroxide is unknown but can be regarded as the hypothetical acid from which stannates, with a formula Sn(OH)62−, are derived by reaction with the (Lewis) basic hydroxide ion.Greenwood, pp. 383–384 Hydrolysis of Pb2+ in aqueous solution is accompanied by the formation of various hydroxo-containing complexes, some of which are insoluble. The basic hydroxo complex Pb6O(OH)64+ is a cluster of six lead centres with metal-metal bonds surrounding a central oxide ion. The six hydroxide groups lie on the faces of the two external Pb4 tetrahedra. In strongly alkaline solutions soluble plumbate ions are formed, including Pb(OH)62−.Greenwood, p. 395 Other main-group elements In the higher oxidation states of the elements in groups 5, 6 and 7 there are oxoacids in which the central atom is attached to oxide ions and hydroxide ions. Examples include phosphoric acid H3PO4, and sulfuric acid O4. In these compounds one or more hydroxide groups can dissociate with the liberation of hydrogen cations as in a standard Brønsted–Lowry acid. Many oxoacids of sulfur are known and all feature OH groups that can dissociate.Greenwood, p. 705 Telluric acid is often written with the formula H2TeO4·2H2O but is better described structurally as Te(OH)6.Greenwood, p. 781 Ortho-periodic acidThe name is not derived from "period", but from "iodine": per-iodic acid (compare iodic acid, perchloric acid), and it is thus pronounced per-iodic , and not as . can lose all its protons, eventually forming the periodate ion IO4−. It can also be protonated in strongly acidic conditions to give the octahedral ion I(OH)6+, completing the isoelectronic series, E(OH)6z, E = Sn, Sb, Te, I; z = −2, −1, 0, +1. Other acids of iodine(VII) that contain hydroxide groups are known, in particular in salts such as the meso''periodate ion that occurs in K4I2O8(OH)2·8H2O.Greenwood, pp. 873–874 As is common outside of the alkali metals, hydroxides of the elements in lower oxidation states are complicated. For example, phosphorous acid H3PO3 predominantly has the structure OP(H)(OH)2, in equilibrium with a small amount of P(OH)3. Holleman, pp. 711–718 The oxoacids of , and have the formula O(''n−1)/2A(OH) where n'' is the oxidation number: +1, +3 or +5, and A = Cl, Br or I. The only oxoacid of fluorine is F(OH). When these acids are neutralized the hydrogen atom is removed from the hydroxide group.Greenwood, p. 853 Transition and post-transition metals The hydroxides of the transition metals and post-transition metals usually have the metal in the +2 (M = Mn, Fe, Co, Ni, Cu, Zn) or +3 (M = Fe, Ru, Rh, Ir) oxidation state. None are soluble in water, and many are poorly defined. One complicating feature of the hydroxides is their tendency to undergo further condensation to the oxides, a process called olation. Hydroxides of metals in the +1 oxidation state are also poorly defined or unstable. For example, silver hydroxide Ag(OH) decomposes spontaneously to the oxide (Ag2O). Copper(I) and gold(I) hydroxides are also unstable, although stable adducts of CuOH and AuOH are known. The polymeric compounds M(OH)2 and M(OH)3 are in general prepared by increasing the pH of an aqueous solutions of the corresponding metal cations until the hydroxide precipitates out of solution. On the converse, the hydroxides dissolve in acidic solution. Zinc hydroxide Zn(OH)2 is amphoteric, forming the zincate ion Zn(OH)42− in strongly alkaline solution. Numerous mixed ligand complexes of these metals with the hydroxide ion exist. In fact these are in general better defined than the simpler derivatives. Many can be made by deprotonation of the corresponding metal aquo complex. :''L''nM(OH2) + ''B L''nM(OH) + ''B''H+ (''L = ligand, B = base) Vanadic acid H3VO4 shows similarities with phosphoric acid H3PO4 though it has a much more complex vanadate oxoanion chemistry. Chromic acid H2CrO4, has similarities with sulfuric acid H2SO4; for example, both form acid salts A+HMO4−. Some metals, e.g. V, Cr, Nb, Ta, Mo, W, tend to exist in high oxidation states. Rather than forming hydroxides in aqueous solution, they convert to oxo clusters by the process of olation, forming polyoxometalates.Juan J. Borrás-Almenar, Eugenio Coronado, Achim Müller Polyoxometalate Molecular Science, Springer, 2003, ISBN 1-4020-1242-X, p. 4 Basic salts containing hydroxide In some cases the products of partial hydrolysis of metal ion, described above, can be found in crystalline compounds. A striking example is found with zirconium(IV). Because of the high oxidation state, salts of Zr4+ are extensively hydrolyzed in water even at low pH. The compound originally formulated as ZrOCl2·8H2O was found to be the chloride salt of a tetrameric cation Zr4(OH)8(H2O)168+ in which there is a square of Zr4+ ions with two hydroxide groups bridging between Zr atoms on each side of the square and with four water molecules attached to each Zr atom.Wells, p. 561 The mineral malachite is a typical example of a basic carbonate. The formula, Cu2CO3(OH)2 shows that it is half-way between copper carbonate and copper hydroxide. Indeed, in the past the formula was written as CuCO3·Cu(OH)2. The crystal structure is made up of copper, carbonate and hydroxide ions. The mineral atacamite is an example of a basic chloride. It has the formula, Cu2Cl(OH)3. In this case the composition is nearer to that of the hydroxide than that of the chloride CuCl2·3Cu(OH)2.Wells, p. 393 Copper forms hydroxy phosphate (libethenite), arsenate (olivenite), sulfate (brochantite) and nitrate compounds. White lead is a basic lead carbonate, (PbCO3)2·Pb(OH)2, which has been used as a white pigment because of its opaque quality, though its use is now restricted because it can be a source for lead poisoning. Structural chemistry The hydroxide ion appears to rotate freely in crystals of the heavier alkali metal hydroxides at higher temperatures so as to present itself as a spherical ion, with an effective ionic radius of about 153 pm. Thus, the high-temperature forms of KOH and NaOH have the sodium chloride structure,Victoria M. Nield, David A. Keen Diffuse neutron scattering from crystalline materials, Oxford University Press, 2001 ISBN 0-19-851790-4, p. 276 which gradually freezes in a monocinically distorted sodium chloride structure at temperatures below about 300 °C. The OH groups still rotate even at room temperature around their symmetry axes and, therefore, cannot be detected by X-ray diffraction. The room-temperature form of NaOH has the thallium iodide structure. LiOH, however, has a layered structure, made up of tetrahedral Li(OH)4 and (OH)Li4 units.Wells, p. 548 This is consistent with the weakly basic character of LiOH in solution, indicating that the Li-OH bond has much covalent character. The hydroxide ion displays cylindrical symmetry in hydroxides of divalent metals Ca, Cd, Mn, Fe, and Co. For example, magnesium hydroxide Mg(OH)2 (brucite) crystallizes with the cadmium iodide layer structure, with a kind of close-packing of magnesium and hydroxide ions. The amphoteric hydroxide Al(OH)3 has four major crystalline forms: gibbsite (most stable), bayerite, nordstrandite and doyleite.Crystal structures are illustrated at Web mineral: Gibbsite, Bayerite, Norstrandite and Doyleite All these polymorphs are built up of double layers of hydroxide ions – the aluminium atoms on two-thirds of the octahedral holes between the two layers – and differ only in the stacking sequence of the layers.Athanasios K. Karamalidis, David A. Dzombak Surface Complexation Modeling: Gibbsite, John Wiley and Sons, 2010 ISBN 0-470-58768-7 pp. 15 ff The structures are similar to the brucite structure. However, whereas the brucite structure can be described as a close-packed structure in gibbsite the OH groups on the underside of one layer rest on the groups of the layer below. This arrangement led to the suggestion that there are directional bonds between OH groups in adjacent layers. This is an unusual form of hydrogen bonding since the two hydroxide ion involved would be expected to point away from each other. The hydrogen atoms have been located by neutron diffraction experiments on αAlO(OH) (diaspore). The O-H-O distance is very short, at 265 pm; the hydrogen is not equidistant between the oxygen atoms and the short OH bond makes an angle of 12° with the O-O line.Wells, p. 557 A similar type of hydrogen bond has been proposed for other amphoteric hydroxides, including Be(OH)2, Zn(OH)2 and Fe(OH)3 A number of mixed hydroxides are known with stoichiometry A3MIII(OH)6, A2MIV(OH)6 and AMV(OH)6. As the formula suggests these substances contain M(OH)6 octahedral structural units.Wells, p. 555 Layered double hydroxides may be represented by the formula Mz+1−xM3+x(OH)2q+(Xn−)q/n·''y''H2O. Most commonly, z = 2, and M2+ = Ca2+, Mg2+, Mn2+, Fe2+, Co2+, Ni2+, Cu2+ or Zn2+; hence q = x. In organic reactions Potassium hydroxide and sodium hydroxide are two well-known reagents in organic chemistry. Base catalysis The hydroxide ion may act as a base catalyst. The base abstracts a proton from a weak acid to give an intermediate that goes on to react with another reagent. Common substrates for proton abstraction are alcohols, phenols, amines and carbon acids. The pKa value for dissociation of a C–H bond is extremely high, but the pKa alpha hydrogens of a carbonyl compound are about 3 log units lower. Typical pKa values are 16.7 for acetaldehyde and 19 for acetone.Ouellette, R.J. and Rawn, J.D. “Organic Chemistry” 1st Ed. Prentice-Hall, Inc., 1996: New Jersey. ISBN 0-02-390171-3. Dissociation can occur in the presence of a suitable base. :RC(O)CH2R′ + B RC(O)CH−R′ + BH+ The base should have a pKa value not less than about 4 log units smaller or the equilibrium will lie almost completely to the left. The hydroxide ion by itself is not a strong enough base, but it can be converted in one by adding sodium hydroxide to ethanol :OH− + EtOH EtO− + H2O to produce the ethoxide ion. The pKa for self-dissociation of ethanol is about 16 so the alkoxide ion is a strong enough base The addition of an alcohol to an aldehyde to form a hemiacetal is an example of a reaction that can be catalyzed by the presence of hydroxide. Hydroxide can also act as a Lewis-base catalyst. As a nucleophilic reagent (Nu) and leaving group (L)]] The hydroxide ion is intermediate in nucleophilicity between the fluoride ion F−, and the amide ion NH2−. pdf The hydrolysis of an ester :R1C(O)OR2 + H2O R1C(O)OH + HOR2 also known as saponification is an example of a nucleophilic acyl substitution with the hydroxide ion acting as a nucleophile. In this case the leaving group is an alkoxide ion, which immediately removes a proton from a water molecule to form an alcohol. In the manufacture of soap, sodium chloride is added to salt out the sodium salt of the carboxylic acid; this is an example of the application of the common-ion effect. Other cases where hydroxide can act as a nucleophilic reagent are amide hydrolysis, the Cannizzaro reaction, nucleophilic aliphatic substitution, nucleophilic aromatic substitution and in elimination reactions. The reaction medium for KOH and NaOH is usually water but with a phase-transfer catalyst the hydroxide anion can be shuttled into an organic solvent as well, for example in the generation of dichlorocarbene. Notes References Bibliography * * * * * *